The Electrochemistry is defined as a branch of chemistry which deals with the properties and behavior of electrolytes in solution and the inter-conversion of chemical and electrical energies.
Derivation of Nernst equation for electrode potential
The Nernst Equation is used to determine the single electrode potential of the cell. The Nernst Equation is derived from emf and Gibbs energy under non-standard conditions. Under standard conditions, the Gibbs free energy equation is then,
ΔGº = – nFE°.
Since, ΔG = AGº + RT In K… (1)
Substituting ΔG = – nFE and ΔG° = – nFE° into equation (1) we have,
– nFE = – nFE° + RT In K
Divide both sides of the equation by – nF, we have,
E = E° – RT/nF In K …(2)
Equation (2) can be rewritten in the form of log base 10, When In is converted to log by 2.303 then,
E = E° – 2.303 RT/nF log K… (3)
The R is 8.314, F is 96,500 coulombs which is equal to 0.0592 V at standard temperature T = 298K, so the equation (3) turns into:
E = E° – 0.0592/n log K.
The equation above indicates that the electrical potential of a cell depends upon the reaction quotient K of the reaction. Thus the concentration of reactants decreases as the redox reaction proceeds where the reactants are consumed. Conversely the products concentration increases due to the increased products formation. During this the cell potential gradually decreases until the reaction is at equilibrium at which ΔG = 0.
Significance:
1. The EMF of the cell can be calculated by:
Zn(s) | Zn2+ (0.024 M) || Zn2+ (2.4 M) | Zn(s).
Understandably the Z2+ ions try to move from the concentrated half cell to a dilute solution where the driving force gives rise to 0.0592 V.
At equilibrium concentrations, the two half cells will have to be equal. In this case, the voltage will be zero.
2. The equilibrium constant K may be calculated using standard cell potential Eº for the reaction.
3. The solubility product from the standard cell potentials can be calculated for the following reaction:
AgCl = Ag+ + Cl–1.
Problem
The emf of a cell measured by means of a hydrogen electrode against a saturated calomel electrode is at 298 K is 0.4188 V. If the pressure of the H2(g) was maintained at 1 atm. Let us calculate the pH of the unknown solution given the potential of reference calomel electrode is 0.2415 V.
Solution:
Given:
Emf of Saturated calomel electrode=0.4188 V
Temperature of saturated calomel electrode=298 K
Pressure of the H2(g)= 1 atm
Potential of reference calomel electrode= 0.2415 V
Formula to be used:
Nernst Equation:
Consider the following redox reaction,
Mn + ne ⇌ M.
For such a redox reversible reaction the free energy change (ΔG) and its equilibrium constant (K) are interrelated as,
Where,
ΔG° = Standard free energy change. The equation (1) is termed as Van’tHoff isotherm.
It will produce electrical energy when there is a decrease in free energy (–ΔG) in the above reaction. In the cell if the reaction involves transfer of ‘n’ number of electrons the ‘n’ Faraday of electricity will flow. When E is the emf of the cell, then the total electrical energy (nEF) produced in the cell is,
–ΔG = nEF (or)
–ΔG° = nEF…(2)
where,
–ΔG = Decrease in free energy change.
–ΔG° = Decrease in standard free energy change.
Comparing equation (1) and (2) it becomes,
T = 298 k
Similarly for oxidation potential
(6)
The above equation (5) and (6) are said to be “Nernst equation for single electrode potential”.
1.2 Reference electrodes: Introduction, construction, working and applications of calomel and Ag / AgCl electrodes.
It is a tendency of a metal to loose or gain electron when it is immersed in its own solution at 25°C and 1M Concentration.
Example: Zn/ZnS04 (1M) at 25°C.
When we build a cell to measure a potential of a chemical reaction, where each pole of the cell shall contain an electrode that transfers electrons from the solution to the wire or vice versa.
There are two main types of electrodes:
Indicator electrodes – Electrodes that respond to an analyte concentration.
An indicator electrode may be as simple as a piece of Pt wire that is inert, so it does not take part in a reaction but simply provides a path for the electrons to flow.
Mostly , an indicator electrode may be made of a metal that takes part in the ½ reaction so indicator electrodes can be simple or more complicated.
Reference electrodes – Electrodes that maintain a fixed potential.
It is an electrode which has a known electrode potential and is stable. The high stability is achieved by employing the redox system, that must contain saturated concentrations in each of the participating solutions of the reaction. The reference electrodes can be used in many ways, but the most important of all is in the electrochemical cell. Here it is used as a half cell in the electrochemical cell to allow for the determination of the other half’s cell potential. It is also used in electrochemical measurements and devices like the DPV and electrochemical biosensor, respectively. The reference electrodes can be classified as aqueous, calomel, non-aqueous and own-constructing.
A best example is the Standard Hydrogen Electrode. This is a great reference electrode because we know its potential is 0 so there is no built in potential that we have to compensate for.
The standard hydrogen electrode, or SHE, is composed of an inert solid like the platinum on which the hydrogen gas is adsorbed, immersed in a solution containing hydrogen ions at unit activity. The half-cell reaction for the SHE is given by ,
2H+ (aq) + 2 e- H2 (g)
and the half-cell potential arbitrarily assigned a value of zero (E0 = 0.000 V). The practical application of SHE is limited by the difficulty in preparing and maintaining the electrode, which is primarily due to the requirement for H2 (g) in the half-cell. Most of the potentiometric methods will employ one of the two other common reference half-cells are the saturated calomel electrode (SCE) or the silver-silver chloride electrode (Ag/AgCl).
These two are widely used as a alternative electrode that are much safer to use.
Silver-Silver Chloride electrode:
This is a secondary reference electrode that issued in analytical instrumentation. It consists of a silver electrode which is coated with AgCl and it is immersed in saturated KCl solution.
The Ag/AgCl electrode is a metal metal salt electrode which consists of narrow glass tube at the bottom of which agar is placed above which saturated solution of KCl is placed and the silver wire is used for electrical connections where it is coated electrolytically with AgCl. The electrode acts as both anode and cathode that depends on the other electrode used.
The electrode can be represented as:
Ag / AgCl (s)/saturated KCl-
Ag: AgCl electrode can act either as anode or cathode depending on the other electrode used.
Silver-Silver Chloride electrode
Anode: Ag + Cl- AgCl + e-
Cathode: AgCl + e- Ag + Cl-
Cell reaction: AgCl + e- Ag + Cl-
The electrode potential may be represented by the Nernst equation as:
E = at 298K
The potential of the Ag-AgCl electrode depends on the concentration of KCl used.
For 0.1N KCl E= 0.290V
1N KCl E= 0.223V
Saturated KCl E = 0.199 V
The electrode potential shall decrease with the increase in the concentration of chloride ions.
The potential of this electrode is based on the ½ reaction and it has potentials of +.222 V if we use 1M KCL or .197 V is used for saturated KCL. Since both the Ag and AgCl are solids the only thing that will vary voltage of this cell is Cl hence the two different potentials. The saturated KCL is more reproducible in the lab because we just add KCl to the electrode until we see crystals and we don’t need to get the concentrations exactly as 1M.
Application:
As a secondary reference electrode in place of calomel electrode.
2. While determining whether the potential distribution is uniform or not in ship hulls and old pipelines protected by cathodic protection.
The Silver-Silver Chloride is slightly expensive to make, so there is a less expensive alternative.
The Calomel Electrode :It consists of mercury at the bottom over which a paste of mercury-mercurous chloride is placed and a solution of potassium chloride is then placed over the paste. The electrical contact can be made by sealing a platinum wire in a glass tube . The electrode is connected with the help of the side tube on the left through the salt bridge with the other electrode to make a complete cell.The potential of the calomel electrode will depend upon the concentration of potassium chloride solution. If the potassium chloride solution is saturated, the electrode is termed as saturated calomel electrode (SCE) and if the potassium chloride solution is 1 N, the electrode is termed as the normal calomel electrode (NCE) while for 0.1 N potassium chloride solution, the electrode is referred to as decinormal calomel electrode (DNCE).’
Calomel electrode can act either as anode or cathode depending on the other electrode used.
Calomel electrode
The half-cell is represented as Hg/ Hg2Cl2 (salt)/ KCl (sat)
The potential for this electrode is governed by the equation.
1/2Hg2 Cl 2(s) + e- ——-> Hg(l) + Cl-
Anode: 2 Hg + 2Cl – Hg2Cl2 + 2e-
Cathode: Hg2Cl2 + 2e- 2 Hg + 2Cl-
Net reaction is: Hg2Cl2 (s) + 2e- 2 Hg + 2Cl -
The electrode potential may be represented by the Nernst equation as
E = Eo – 0.059/n log [Cl -]2
E = Eo – 0.0591/2 log [Cl-]2
E = Eo – 0.0591 log [Cl ] at 298 K
The potential of the calomel electrode will depend on the concentration of KCl used.
For 0.1N KCl E = 0.33V
1N KCl E = 0.28V
Saturated KCl E = 0.24V
The electrode potential decreases with increase in the concentration of chloride ions.
The electrode potential of any other electrode on the hydrogen scale can be measured when it is combined with calomel electrode. The emf of such a cell is measured. Thus, from the value of electrode potential of calomel electrode, the electrode potential of the other electrode can be evaluated.
Again all but Cl are solids to we get 2 potentials +.268 if KCl=1M +.241 if KCl is saturated.
Given a potential measured against any one of these three common reference electrodes we should be able to convert the potential to measure against one of the other electrodes.
For example, we measured a new potential gains of an SHE electrode and found it has a +.2V potential then the potential we measured would be against a Saturated Calomel Electrode.
Our measured potential is +.2 relative to 0 and the SHE is at +.241 from 0, so using this electrode we would measure a +.041 (.241-.2) potential.
We measure a potential of +.1 relative to a SHE electrode then the measure relative to a Ag/AgCl electrode is
.241 + .1 = .341
.341-.197 = .144.
Advantages of calomel reference electrodes:
Easy to construct
2. Easy to maintain.
3. The cell potential is stable over a long period and does not vary with temperature
4. Oxidizing agents can also be used.
1.2.1 Measurement of electrode potential using Calomel electrode (CE)
To measure the potential of given electrode using calomel electrode, a galvanic cell is obtained by combining both the electrodes according to cell conventions.
If the given electrode undergoes oxidation with respect to CE, then it is taken as anode and CE is taken as cathode.
Measurement of electrode potential using calomel electrode
The galvanic cell obtained in this case can be represented by:
M | M+n || KCl | Hg2Cl2(s) | Hg(l)
At anode, oxidation occurs, M M+n + ne-
At cathode, reduction occurs, Hg2Cl2 + 2e- 2 Hg + 2Cl-
EMF of cell, ECell = Ecathode – EAnode
ECell = ECE – EMn+/M
EMn+/M = ECE – Ecell
If the given electrode undergoes reduction with respect to CE, then it is taken as cathode and CE is taken as anode.
Measurement of electrode potential using calomel electrode
The galvanic cell obtained in this case can be represented by:
Hg(l) | Hg2Cl2(s) | KCl || M+n |M
At anode, oxidation occurs, 2 Hg + 2Cl- Hg2Cl2 + 2e-
At cathode, reduction occurs, M+n + ne- M
EMF of cell, ECell = Ecathode – EAnode
ECell = EMn+/M__─ ECE
EMn+/M = ECE + Ecell
1.3 Ion selective electrode: Introduction
Ion-selective electrode
It possesses the ability to respond only to certain specific ions thereby developing a potential with respect to that species only in a mixture and ignoring the other ions totally. In other words the potential developed by an ion selective electrode depends only on the concentration of species or ions of interest.
For example, the glass membrane is only H+ ions selective. The material employed for some ion selective membranes are mentioned below:
Solid state membranes:
(i) For fluoride ions: lanthanum tri fluoride crystal doped with europium di fluoride.
(ii) For chloride ions: pressed pallet of Ag2S + AgCl. The electrode has a teflon body and the pellet is held in position by using epoxy resin.
Liquid state membranes:
(i) For certain alkali and alkaline earth cations: neutral monocyclic crown ethers and phosphatediester.
(ii) For anions: iron phenanthrolic complexes.
The liquid state membrane is usually obtained by absorbing the active molecules on an inert porous support such as porous polymer.
Ion selective electrode
Limitations:
(i) Usually ion selective membranes can be used in solution of pH value upto 10, since higher pH values affect these.
(ii) Although ion selective membranes are very thin, yet their resistances are extremely high, so it is necessary to use electronic potentiometers to measure the potential difference.
Applications:
In the determination of:
(i) Concentration of cations, e.g., H+, Li+, Na+, K+, NH4+, Ag+, Pb2+, Cd2+, hardness and anions, e.g., halide ions(X-), NO3-, CN-, S2-etc.
(ii) pH of a solution by using H+ion-selective glass electrode.
(iii) concentration of a gas by using gas–sensing electrode. For example, glass electrode is employed for measuring the CO2 (g) level in blood. The glass electrode is kept in contact with a very thin-walled CO2(g) permeable silicone rubber membrane soaked in dilute NaHCO3 solution.
Then the electrode is dipped in the blood sample under-test. Since the membrane is permeable to CO2 (g), So CO2 (g)permeates into the membrane thereby causing reaction between NaHCO3 and CO2 (g).
Eventually the pH value is altered, which is sensed by the glass electrode. In other words observed potential of such a glass electrode can be employed to measure the concentration of CO2 (g) in the blood sample.
1.3.1 Construction and working of glass electrode, determination of pH using glass electrode
Glass electrode
When the two solutions of different pH values are separated by a thin glass membrane, then it develops a difference of potential between the two surface of the membrane. The potential difference developed is proportional to the difference in pH value.
The glass membrane functions as the ion-exchange resin and an equilibrium is set up between the Na+ ions of glass and H+ ions in solution.
For a particular type of glass the potential difference varies with the H+ ion concentration, and is given by the expression:
EG= E0G+ 0.0592 V pH
Over a range of pH of the test solutions from 0 to 10.
Construction:
A glass electrode consists of thin-walled glass bulb containing AgCl coated Ag electrode or simply a Platinum electrode in 0.1 M –HCl. The glass electrode may be shown schematically as:
-Ag | AgCl(s), HCl(0.1M) | Glass or Pt, 0.1 M HCl | Glass+
HCl in the bulb furnishes a Constant H+ ion concentration. Thus it is a silver-silver chloride electrode,reversible with respect of chloride ions. The glass electrode is used as “internal reference electrode”.
For the determination of pH of solution, especially colored solution containing oxidizing or reducing agents. Usually calomel electrode is used as the second electrode. In order to determine the pH of a solution, it is placed is the solution under –test and this half-cell is coupled with saturated calomel electrode.
The e.m.f. of the cell is measured. Due to the higher resistance, special electron-tube voltmeters are used to measure the e.m.f of the above cell.
Determination of pH using glass electrode
In the glass-electrode method, the known pH of a reference solution is determined by using the two electrodes, a glass electrode and a reference electrode, and also by measuring the voltage generated between the two electrodes. The difference in pH between the solutions inside and outside the thin glass membrane creates electromotive force in proportion to this difference in pH. Hence, this thin membrane is known as the electrode membrane. Normally, if the temperature of the solution is 30 °C, when the pH inside is different from that of outside by 1, it can create approximately 60 mV of electromotive force.
Liquid inside the glass electrode usually has a pH of 7. Thus, if one measures the electromotive force which is generated at the electrode membrane, the pH of the sample will be found by calculation.
A second electrode is necessary while measuring the electromotive force generated at the electrode membrane of a glass electrode. The other electrode, paired with the glass electrode, is known as the reference electrode. The reference electrode must have extremely stable potential. Thus, it is provided with a pin hole or a ceramic material at the liquid junction.
It can also be explained as, a glass electrode is devised to generate accurate electromotive force due to the difference in pH and also a reference electrode is devised not to cause electromotive force due to a difference in pH. The two electrodes are connected to potentiometer or pH meter. Thus, the combined electrodes can be represented as:
Hg(l) / Hg2Cl2(S) / Saturated KCl //solution of unknown pH /glass/0.1M Hcl/Ag/AgCl(s)
The e.m.f. of the complete cell is given by,
The potential of E glass is given by:
Where a1= Concentration of [H+] ions in glass bulb
A2= Concentration of [H+] ions in unknown solution
Eglass = RT/nf ln a1-RT/nf ln [a2]
Eglass = Constant-RT/nf ln [a2] ( since [H+]ions in glass bulb is constant)
Eglass = Constant – 0.0591 log[H+]
Eglass = Constant – 0.0591 pH
= 0.2422 V –[ E0G+ 0.0592 V pH]
Advantages of glass electrode:
It is simple and can easily be used.
Equilibrium is rapidly achieved.
The results are accurate.
It is easily poisoned.
Limitations:
(1) It can be used in solutions with pH range of 0 to 10. The electrodes composed of special glasses can be used for measurements upto a pH of 12. But the pH cations of solution above 12 will affect the glass interface and render the electrode useless.
(2) Although glass membrane of electrode is very thin yet its resistance is extremely high, that cannot be measured by ordinary potentiometers. It is therefore necessary to use special electronic potentiometers.
1.3.2 Concentration cells: Electrolyte concentration cells, numerical problems
Concentration Cells:
It is a type of galvanic cell. When the electrode and electrolyte present in both the half cells are same but only the concentration of metal or electrolyte is different is known as the concentration cell .
Example:
Consider the following concentration cell that is constructed by dipping two copper electrodes in CuSO4 solutions of M2 molar and M1molar where M2M > M1M. The two half-cell are internally connected by a salt bridge and externally connected by a metallic wire through voltmeter or ammeter.
Concentration cell
The electrode, which is dipped in less ionic concentrations solutions (M1) act as anode and undergoes oxidation. The electrode, which is dipped in more ionic concentration (M2) act as cathode and undergoes reduction.
Types of concentration cells:
There are two types of concentration cells:
1. Electrode concentration cell
2. Electrolyte concentration cell .
Electrode concentration cell:
In electrode concentration cell, the potential difference is developed between two similar electrodes at different concentrations which is dipped in the same electrolytic solution.
Example: Two hydrogen electrodes at different gaseous pressure that are dipped in same electrolyte solution.
Cell representation:
Pt, H2 (P1) / H2 (P2), Pt
If P1>P2 oxidation occurs at L.H.S. The electrode and reduction occurs at R.H.S.
Cell reaction is
L.H.S. H2 (P1) 2H++2e
R.H.S. 2H+ + 2e – H2 (P2)
Overall cell reaction is H2 (P1) H2 (P2)
Electrolyte concentration cells:
It is a type of galvanic cell in which the electrode & electrolyte present in both the half cells are same but only the concentration of the electrolyte is different.
Construction of an electrolyte concentration cell
Cell representation:
M / M n+(C1) // M n+(C2) / M
The metal immersed in the dilute solution will have lower potential act as anode, is shown as:
M Mn+ (C1) + ne-
The metal immersed in the concentrated solution will have higher potential act as cathode, is shown as:
Mn+ (C2) + ne- M
The Net cell reaction is:
Mn+ (C2) Mn+ (C1)
The Nernst equation for electrolyte concentration cell is:
ECell = EoCell + 0.0591 log [Mn+]cathode
n [Mn+] Anode
EoCell =Eocathode – Eoanode = 0
ECell = 0.0591 log [Mn+]cathode
n [Mn+] Anode
ECell = 0.0591 log C2 at 298k
n C1
ECell is +ve and reaction is spontaneous only when C2 > C1.
Applications:
1. To determine the valence of ions.
2. To determine the solubility product of sparingly soluble salts.
3. To determine the equilibrium constant.
Problems:
1. Let us write the electrode reactions and Calculate the EMF of the given cell at 298K, Ag(s)AgNO3 (0.018M) AgNO3 (1.2M)Ag(s).
Solution:
Given:
Ag(s)AgNO3 (0.018M) AgNO3 (1.2M)Ag(s)
Temperature= 298K
Formula to be used:
At anode: Ag(s) Ag+ + e-
At Cathode: Ag+ + e- Ag(s)
w.k.t at 298K
(n=1)
Ecell = 0.1078 V.
2. Let us calculate the emf of Copper concentration cell at 250 C, where the copper ions ratio in the cell is 10.
Solution:
Given:
Temperature= 250 C
Ratio of copper ions= 10
Formula to be used:
w.k.t ; at 298 K
Ecell = 0.0296 V.
3. A cell contains two hydrogen electrodes. The negative electrode is in contact with a solution of 10-6 M hydrogen ions. The emf of the cell is 0.118volt at 25° C. Calculate the concentration of hydrogen ions at the positive electrode.
Solution:
Given:
Emf = 0.118V
Temperature= 250 C
Negative electrode of hydrogen ions= 10-6 M
Formula to be used:
The cell may be represented as:
Pt|H2(1 atm)|H+||H+|H2(1 atm)|Pt
10-6 M
Anode: H2 → 2H+ + 2e-
Cathode: 2H+ + 2e → H2
Ecell = 0.0591/2 log([H+ ]cathode2)/[10-6 ]2
0.081 = (0.0591) log ([H+])/10-6
log[H+ ]cathode/10-6 =0.118/0.0591=2
[H+ ]cathode/10-6 = 102
[H+]cathode = 10-6 = 10-4 M
4. The emf of the cell Ag|Agl in 0.05 MK\\Sol. NH4NO3|10.05 M AgNO3\\Ag is 0.788 volt at 25°C. The activity coefficient of KI and silver nitrate in the above solution is 0.90 each. Calculate (i) the solubility product of Agl, and (ii) the solubility of Agl in pure water at 25°C.
Solution:
Given:
Emf of Ag|Agl = 0.05 MK\\Sol
Emf of NH4NO3 = 10.05 M
Emf ofAgNO3\\Ag = 0.788 V
Temperature= 250 C
Activity coefficient of KI and silver nitrate= 0.90
Formula to be used:
Solubility of Agl = √(Solubility product of Agl)
Ag+ ion concentration on AgN03 side = 0.9 × 0.5 = 0.045 M.
Similarly the I- ion concentration in 0.05 M KI solution = 0.05 × 0.9 – 0.045 M.
Ecell = 0.0591/1 log[Ag+ ](R.H.S.)/[Ag+ ](L.H.S.)
= 0.0591 log 0.045/[Ag+ ](L.H.S.)
or
log 0.045/[Ag+ ](L.H.S.) = 0.788/0.0591 = 13.33
[Ag+]L.H.S. = 0.045/(2.138× 1013 ) = 2.105 × 10-15 M
Solubility product of Agl = [Ag+][I-]
= 2.105 × 10-15 × 0.045
= 9.427 × 10-17.
Solubility of Agl = √(Solubility product of Agl)
Solubility of Agl = √(9.472×10^(-17) )
= 9.732 × 10-9 g mol L-1
= 9.732 × 10-9 × 143.5 g L-1
= 1.396 × 10-6 g L-1.
1.4 Battery Technology: Introduction, classification – primary, secondary and reserve batteries.
Battery:
A battery is a device. It consists of two or more galvanic cell that are connected in series or parallel or both, which converts the chemical energy into electrical energy through the redox reaction.
Some examples for battery are Lithium ion battery, Lead acid battery, Nickel-Cadmium battery etc.
Applications:
It is used in calculators, watches, emergency lightning in hospitals, car engines, space applications, military, computers and also in pacemakers for heart hearing aids.
Basic Components of Battery:
It consists of four major components. They are:
Anode :
It is a negative electrode which releases the electrons into the external circuit by undergoing oxidation.
M Mn+ + e-
The above equation represents the anode reaction which takes place in the circuit.
Cathode :
It is a positive electrode which accepts the electrons coming from anode through external circuit.
Mn+ + e- M
The above equation represents the cathode reaction which takes place in the circuit.
Electrolyte:
The electrolyte provides a medium for transferring the ions inside the cell between the anode and cathode. Any solution of an acid, alkali or salt which has high ionic conductivity can commonly be used as an electrolyte.
Separator:
The separator is used to separate the anode and the cathode compartments in a battery for preventing the internal short circuiting. It allows the ions from anode and cathode. For example, the Cellulose and nafion membranes.
Classification Of Batteries:
The batteries are classified as:
Primary battery or primary cells: The primary battery is said to be a battery that cannot be recharged, cell reactions are irreversible and it can be discarded when the battery has delivered all its electrical energy.
Examples: dry cell or Li-Mno2 cell , Zn-MnO2 Cell.
Secondary battery: The secondary battery is said to be a battery that can be recharged, cell reactions are reversible and after discharging, it can be recharged.
Examples: Nickel cadmium cell and lead storage cell .
Reserve Batteries: In the reserve batteries, one of the components is stored separately and is incorporated into the battery when required.
Examples: Mg-CuCl and Mg-AgCl battery.
1.4.1 Characteristics – cell potential, current, capacity, electricity storage density, energy efficiency, cycle life and shelf life
Characteristics of a battery:
Voltage or cell potential
Current
Capacity
Electricity storage density
Energy Efficiency
Cycle Life
Self Life
1. Voltage or cell potential (EMF):
The battery voltage is given by the equation:
Ecel l= (EC – EA) -ηA – ηC – iRcell
Where,
EC and EA are the Electrode potentials of cathode and anode
ηA and ηC are the over potentials at cathode and anode
iRcell is the internal resistance of the cell.
The conditions to derive maximum voltage from a battery:
1. The potential difference must be high.
2. Internal resistance of the cell must be low.
3. Over potentials at cathode and anode should be minimum.
2. Current:
For a efficient discharge, the electron must flow at a uniform rate in the electrolyte current is a measure of the rate of flow of electrons during discharge . Current is amount of charge flowing per unit and it it is expressed in ampere per second. The batteries provide direct current.
According to Ohms law, current I=V/R . In order to force the current through the cell, more potential difference is required at higher resistance.
Capacity:
The battery capacity is a measure of the charge which is stored by a battery. The corresponding SI unit is Ampere-hour (Ah).The theoretical capacity may be calculated using faradays relation,
C = wnF/M,
Where, C is the capacity
w is the mass of the active material
M is its molar mass
n is the number of moles electrons.
It is the amount of electrical energy the battery that delivers over certain period. Capacity also depends upon the size of the battery.
Electricity storage density:
Electricity storage density is a measure of the charge per unit mass which is stored in the battery. The mass of battery includes terminals, the masses of electrolyte current collectors and other subsidiary elements. The lighter subsidiary elements leads to high storage density.
Example: A 7g of lithium anode gives 96500 C whereas, for the same charge, 65g of zinc would be required.
Energy efficiency:
The energy efficiency is defined as the ratio of useful energy output to the total energy input. A battery must have high energy efficiency.
It can be represented as,
Cycle-life:
The cycle life of a battery is defined as the number of charge-discharge cycle which can be achieved before the failure occurs. If the average depth of discharge is greater, the cycle life will be shorter. This can be used only for the secondary batteries.
It is necessary that during charging the active material is regenerated in a suitable state for the discharge. Therefore, the discharge-charge cycle depends on chemical composition and distribution of active materials in the cell.
Shelf-life:
Certain batteries can be stored for many years. Therefore, the duration of storage of a cell without the self discharge or the loss of performance is called shelf-life.
1.4.2 Construction, working and applications of Zinc- Air, Nickel- metal hydride batteries
Zinc-Air Battery
Anode: Granulated Zn powder
Cathode: Air/C
Electrolyte: KOH 6M
The production of electrochemical energy in Zn/air battery is due to the use of oxygen from the atmosphere. The diffused oxygen acts as a cathode reactant in the battery. The air cathode catalytically promotes the reaction of oxygen with an aqueous alkaline electrolyte and it does not consume or change during the discharge. When an alkaline electrolyte is used in the Zn/air battery, it is necessary to increase only the amount of zinc present to increase battery capacity.
A typical zinc-air cell
The air cathode acts only as a reaction site and is not consumed. The reason for the increased energy density in the Zn/air cell is because of the larger volume containing the active material. Since the air cathode has infinite life, the electrical capacity of the cell is determined only by the anode capacity, resulting in at least a doubling of energy density.
Theoretically, the air cathode has infinite use life and its physical size and its electrochemical properties remain unchanged during cell discharge. A schematic representation of a typical Zn/air cell is shown in the above figure. A loose granulated powder of Zn mixed with electrolyte [KOH] acts as the zinc anode material and in some cases, a gelling agent is used to mobilise the composite and ensure adequate electrolyte contact with zinc granules. The outer metal (button type) acts as the cathode of the battery and a plastic gasket insulates the anode active materials and the cathode as shown in the above figure.
The chemistry of the electrode reactions taking place in zinc/air battery are as follows:
At anode,
At cathode,
The overall efficiency of zinc/air battery,
The output of the zinc-air battery is 1.65 V.
Advantages
Zn/air battery technology offers the following advantages for many applications.
High energy density
Flat discharge voltage
Long shelf life
No ecological problems
Low cost
Capacity independent of load and temperature.
Applications:
Zn-air batteries have been most successfully employed as a power source for hearing aids. Other applications include electronic pagers, voice transmitters, portable battery chargers, various medical devices and so on.
Nickel-Metal Hydride Cell
(i) Anode — A metal hydride, MH
(ii) Cathode — Nickel oxy hydroxide/Ni
(iii) Electrolyte — KOH.
A relatively new technology is adopted in the case of the chargeable sealed nickel-metal hydride battery with characteristics similar to those of the sealed Ni-Cd batteries. The Ni-MH battery uses hydrogen absorbed in a metal alloy for the active negative material whereas cadmium is used in the Ni-Cd battery and that makes the noticeable difference between the two.
A higher energy density can be achieved in the case of metal hydride electrode than the cadmium electrode. Thus, a smaller amount of the negative electrode is used in the Ni-metal hydride. This allows for a larger volume for the +ve electrode, which results in a higher capacity or longer service life for the metal hydride battery. Moreover, as the metal hydride battery is free of Cd, it is considered more environmentally friendly than the Ni-Cd battery and may reduce the problems associated with the disposal of rechargeable nickel batteries.
Nickel-metal hydride batteries consist of a positive plate of a highly porous sintered or felt nickel substrate impregnated with nickel hydroxide as its principal active material, a negative plate of a highly porous structure using a perforated LaNi5 alloy grid (a hydrogen-absorbing alloy). A synthetic non-woven material separates the two electrodes, which serves as a medium for absorbing the electrolyte and a sealing plate provided with a self-resealing safety vent.
The chemistry of the electrode reactions of Ni-MH battery can be described as follows:
Ni-MH cell
In the charged state of Ni-MH battery, nickel oxy-hydroxide is the active material of the +ve electrode. This is same as the positive electrode in the Ni-Cd battery.
In the charged state of the Ni-MH battery, hydrogen is stored in a hydrogen absorbing alloy as metal hydride, LaNi5, (—ve active material). This metal alloy is capable of undergoing a reversible hydrogen absorbing-desorbing reaction as the battery is charged and discharged.
An aqueous solution of KOH is the major component of the electrolyte, with minimum amount of the electrolyte absorbed by the separator and the electrodes. As can be seen from the overall reaction given below, the chief characteristics of the principle behind a Ni-MH battery is that hydrogen moves from the +ve to —ve electrode during charge and in reverse order during the discharge. with the electrolyte taking no part in the reaction; which means that there is no accompanying increase or decrease in electrolyte.
The discharge electrode reactions of the Ni-MH battery are as follows:
At anode,
At cathode,
The nickel oxyhydroxide is reduced to nickel hydroxide.
The overall reaction on discharge is,
The process is reversed during charge of the Ni-MH battery.
Advantages
The following are the advantages of an Ni-MH battery:
High capacity
No maintenance required
Minimum environmental problem
Rapid recharging capability
Long cycle life
Long shelf life in state of charge.
Applications
Nickel-metal hydride batteries are used in computers, cellular phones and other portable and consumer electronic applications where high specific energy is required.
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